1. History of the Atomic Model

2. Early Greek Theories Aristotle (350 B.C.) 4 Elements
Democritus (400 B.C)
Atoms and a void (empty space)
Atoms are indivisible
Based on philosophical arguments, NOT experimental evidence

3. Dalton’s Billiard Ball Model (1805)
All matter is made of tiny, indivisible particles called atoms.
Each element is made up of atoms that are unique and unlike atoms of any other element.
Atoms of an element are identical.
Matter is composed of indestructible, indivisible atoms

4. Thomson’s Raisin Bun Model (1897)
Materials, when rubbed, can develop a charge difference.
This electricity was called “cathode rays”
These rays have a small mass and are negatively charged.
Thomson noted that these negative subatomic particles (electrons) were a fundamental part of all atoms.

5. Rutherford’s Nuclear Model (1911)
Rutherford shot alpha () particles at gold foil.
Zinc sulfide screen
Thin gold foil
Lead block
Radioactive substance
path of invisible -particles
Most particles passed through. So, atoms are mostly empty space.
Some positive -particles deflected or bounced back!
Thus, a “nucleus” is positive (protons) & holds most of an atom’s mass.

6. Table 1, p. 26

7. Practice
P. 26 #1-5

8. Atomic numbers, Mass numbers
Elements are often symbolized with their mass number (A) and atomic number (Z)
E.g. Oxygen: O 16 8
Z = # of protons = # of electrons
A - Z = # of neutrons
Calculate # of e–, n0, p+ for Ca, Ar, and Br

9. Atomic
Mass p+ e- no Ca 20 40 20 20 20 Ar 18 40 18 18 22 Br 35 80 35 35 45

10. Bohr - Rutherford diagrams
Putting all this together, we get B-R diagrams
To draw them you must know the # of protons, neutrons, and electrons (2,8,8,18 filling order)
Draw protons (p+), (n0) in circle (i.e. “nucleus”)
Draw electrons around in shells
3 p+
4 n0
2e– 1e–
Li shorthand
2 p+
2 n0 He Li
Draw Be, B, Al and shorthand diagrams for O, Na

11. Be B Al O Na 8 p+ 11 p+ 8 n° 12 n° 4 p+ 5 n° 5 p+ 6 n° 13 p+ 14 n°
2e– 8e– 1e– Na
8 p+
8 n°
2e– 6e– O

12. Isotopes and Radioisotopes
Isotopes: Atoms of the same element that have different numbers of neutrons
Due to isotopes, mass #s are not round #s.
E.g. Li (6.9) is made up of both 6Li and 7Li.
Often, at least one isotope is unstable.It breaks down, releasing radioactivity.These types of isotopes are called radioisotopes
Q- Sometimes an isotope is written without its atomic number - e.g. 35S (or S-35). Why?
Q- Draw B-R diagrams for the two Li isotopes.
A- The atomic # of an element doesn’t change Although the number of neutrons can vary, atoms have definite numbers of protons.

13. 6Li 7Li
3 p+
3 n0
2e– 1e–
3 p+
4 n0
2e– 1e–

14. Practice
P. 29 #1-7

15. Limitations to Rutherford’s Model
Orbiting electrons should emit light, losing energy in the process
This energy loss should cause the electrons to collapse into the nucleus
However, matter is very stable, this does not happen

16. Bohr’s Planetary Model
Electrons orbit the nucleus in energy “shells”
An electron can travel indefinitely within a shell without losing energy
The greater the distance between the nucleus and the shell, the greater the energy level
An electron cannot exist between shells, but can move to a higher, unfilled shell if it absorbs a specific quantity of energy, or to a lower, unfilled shell if it loses energy (quantized)
When all the electrons in an atom are in the lowest possible energy levels, it is in its ground state.

17. An atom becomes excited when one of its electrons absorb energy
If enough energy is absorbed then the electron can make a quantum leap to the next energy level, if there is room
When the electron returns to a lower energy state the energy is released in the form of a photon, which we see as visible light

18. The energy of the photon determines its wavelength or color
Each element has its own frequencies of color, so it emits its own distinctive glow

20. Summary of Atomic Models
Dalton’s “Billiard ball” model ( )
Atoms are solid and indivisible.
Thomson’s “Raisin bun” model (1900)
Negative electrons in a positive framework.
Rutherford’s “Nuclear” model (~1910)
Atoms are mostly empty space.
Negative electrons orbit a positive nucleus.
4) Bohr’s “Planetary” model (~1920)
Negative electrons orbit a positive nucleus.
Quantized energy shells
5) Quantum Mechanical model (~1930)
Electron probabilities (orbitals)

21. Practice
Pre-lab: Atomic Spectra (p. 40)
p. 42 #1-3
p. 45 #6-8